Now that we’ve defined acids and bases, let’s discuss
how they work together in reactions. Look at the generic acid-base
reaction below:

HX(aq) + H2O(l) X-(aq) +
H3O(aq)

When the forward reaction occurs, HX donates a proton
to water (so it acts as the base) to form hydronium. When the reverse
reaction occurs, the hydronium ion acts as the acid, donating a
proton to the X-. Together, HX and X- are
said to be conjugate acid-base pairs. Conjugate acid-base pairs
are compounds that differ by the presence of one proton, or H+. All
acids have a conjugate base, which is formed when their proton has
been donated; likewise, all bases have a conjugate acid, formed
after they have accepted a proton.

Example

Apply the appropriate acid-base theory to first identify
the acid and base reacting and then identify the conjugate acid-base
pairs in the examples below:

Explanation

In this first reaction, we see that HNO3 gives
a proton to water, which then forms a hydronium ion. This makes
HNO3 the acid in the forward reaction, and
water acts as the base. HNO3’s conjugate
base is , and water’s conjugate
acid is the hydronium ion, or

Here donates the proton to water, so
in the forward reaction it acts as the acid, and water is still
the base. ’s conjugate base
is NH3, and water’s conjugate acid is again
the hydronium ion, H3O+.

Relative Strengths of Acids and Bases

Certain acids are stronger than other acids, and some
bases are stronger than others. What this means is that some acids
are better at donating a proton, and some bases are better proton
acceptors. A strong acid or base dissociates or ionizes
completely in aqueous solution. A weak acid or base
does not completely ionize.

Strong Acids

There are six strong acids that you’ll need to memorize
for the SAT II Chemistry test:

Hydrohalic acids: HCl, HBr, HI

Nitric acid: HNO3

Sulfuric acid: H2SO4

Perchloric acid: HClO4

Let’s take a closer look at how acids differ in strength
by focusing on perchloric acid. In general, the greater the number
of oxygen atoms in a polyatomic ion, the stronger the acid.

So HClO4 is stronger than HClO3,
which is stronger than HClO2, which is stronger
than HClO. (Perchloric acid is the strongest among the six, but
all the other oxyacids of chlorine are not considered strong acids.)
Now think about why, as you take away oxygens, the strength of the
acid decreases. The hydrogen (proton) to be removed is bonded to
an oxygen atom. The oxygens are highly electronegative and are pulling
the bonded pair of electrons away from the site
where the hydrogen is bonded, thus making it easier to remove the
H+. As the number of oxygen decreases,
the molecule becomes less polar, and the H+ is harder
to remove.

Strong Bases

There are fewer strong bases to memorize for the exam.
These are hydroxides (—OH), oxides of 1A and 2A metals (except Mg
and Be), H-, and . Remember that the stronger the
acid, the weaker its conjugate base, and the converse is also true.
The figure below illustrates the relative strengths of some common
conjugate acid-base pairs.

The pH Scale

As you know, water can act as either a proton donor (in
the form of the hydronium ion, H3O+)
or a proton acceptor (as OH-). In solution,
a water molecule can even donate a proton to or accept a proton
from another water molecule, and this process is called autoionization:

2H2OH3O+ +
OH-

Since this reaction takes place in equilibrium, we can
write an equilibrium expression, Keq, for
it:

Keq =
[H3O+][OH-]

And since this expression refers specifically to the ionization
of water, we can write the equilibrium expression as Kw.
At 25ºC, the value of Kw,
which is known as the ion-product constant, is 110-14.
This means that the [H3O+]
= [OH-] and each is equal to 110-7. When
the concentrations of H+ and OH- are
equal in a solution, the solution is said to be neutral. In acidic
solutions, the concentration of H+ is
higher than that of OH-, and in basic solutions,
the concentration of OH- is greater than
that of H+.

The pH of a solution is calculated as the
negative logarithm in base 10 of the hydronium ion concentration—it
is an expression of the molar concentration of H+ ions
in solution:

pH = -log [H+]
or -log [H3O+]

A solution like the equilibrium expression for water,
which is neutral at standard temperature, would have a pH of

pH = -log [110-7] = -(-7.00)
= 7.00

So as you can see, neutral solutions have a pH of 7. If
the solution contains more hydronium ions than this neutral solution
([H+] > 110-7),
the pH will be less than 7.00, and the solution will be acidic;
if the solution contains more hydroxide ions than this neutral solution
([OH-] > 110-7),
the pH will be greater than 7.00, and the solution will be basic.

Similarly, the pOH of a solution is calculated as the
negative logarithm in base 10 of the hydroxide ion concentration:

pOH = -log [OH-]

and pH and pOH are related to each other by the equation

pH + pOH = 14

Since you won’t be allowed to have a calculator for the
SAT II Chemistry test, you can use the following equation if you
need to calculate the hydronium ion concentration of a solution:

[H3O+]
= 10-pH

Now try a problem: What is the pH of a solution at 25ºC
in which [OH-] = 1.010-5M?

Explanation

The fact that this solution is at 25ºC tells us that we
should use the Kw relationships.
If the [OH-] = 1.010-5 M,
then pOH = 5. You know that 1.010-5 is
the same as plain old 10-5. The log of
10-5 is -5 (simply use the exponent when
a number, any number, is written as 10power,
so the “negative” of the log is equal to -(-5), or simply 5. Now,
if the pOH is 5, then the pH is 9 since pH + pOH = 14.